Doc Brown's A Level Chemistry Advanced Level Theoretical Physical Chemistry – AS A2 Level Revision Notes – Basic Thermodynamics
GCE Thermodynamics–thermochemistry sub–index links below
Part 1 – ΔH Enthalpy Changes – The thermochemistry of enthalpies of reaction, formation, combustion and neutralisation
Part 1.1 Advanced Introduction to Enthalpy (Energy) Changes in Chemical Reactions
This page is an introduction to advanced level ideas on exothermic or endothermic energy changes in chemical reactions – referred to as 'enthalpy changes' e.g. enthalpy of reaction, enthalpy of formation, enthalpy of combustion and also bond enthalpies ('bond energies'). The concept of enthalpy level diagrams is described & introduced and enthalpy changes are clearly distinguished from activation energies. The notion of 'standard conditions' is described and why it is necessary to have data based on a standard defined temperature, pressure and concentration.
Energetics index: GCSE Notes on the basics of chemical energy changes – important to study and know before tackling any of the three Advanced Level Chemistry pages Parts 1–3 here * Part 1a–b ΔH Enthalpy Changes 1.1 Advanced Introduction to enthalpy changes – reaction, formation, combustion : 1.2a & 1.2b(i)–(iii) Thermochemistry – Hess's Law and Enthalpy Calculations – reaction, combustion, formation etc. : 1.2b(iv) Bond Enthalpy Calculations : 1.3a–b Experimental methods for determining enthalpy changes and treatment of results : 1.4 Some enthalpy data patterns : 1.4a The combustion of linear alkanes and linear aliphatic alcohols : 1.4b Some patterns in Bond Enthalpies and Bond Length : 1.4c Enthalpies of Neutralisation : 1.4d Enthalpies of Hydrogenation of unsaturated hydrocarbons and evidence of aromatic ring structure in benzene : Extra Q page A set of practice enthalpy calculations with worked out answers ** Part 2 ΔH Enthalpies of ion hydration, solution, atomisation, lattice energy, electron affinity and the Born–Haber cycle : 2.1a–c What happens when a salt dissolves in water and why? : 2.1d–e Enthalpy cycles involving a salt dissolving : 2.2a–c The Born–Haber Cycle *** Part 3 ΔS Entropy and ΔG Free Energy Changes : 3.1a–g Introduction to Entropy : 3.2 Examples of entropy values and comments * 3.3a ΔS, Entropy and change of state : 3.3b ΔS, Entropy changes and the feasibility of a chemical change : 3.4a–d More on ΔG, Free energy changes, feasibility and applications : 3.5 Calculating Equilibrium Constants : 3.6 Kinetic stability versus thermodynamic feasibility * PLEASE note that delta H/S/G values vary slightly from source to source, so I apologise in advance for any inconsistencies that may arise as I've researched and developed each section.
in what you might call a lower level introduction which bridges GCSE and AS level and you are completely ok in interpreting enthalpy level and activation energy diagrams such as ...
... and can clearly distinguish between enthalpy change and activation energy and the activation energy change with a catalyst does NOT change the enthalpy value of the reaction (activation energy will be rarely mentioned until the end of Part 3), but all important ideas from the GCSE page will be re–studied in their advanced level context, but I would still study the GCSE page first!
REMEMBER – all chemical changes are accompanied by energy changes or energy transfers, many of which can be directly measured, or, theoretically calculated from known values.
1.1b Enthalpy Changes and Thermochemistry
Some important initial definitions and examples:
The system: The reactants and products of the reaction being studied i.e. the contents of the calorimeter.
The surroundings: The means the rest of the 'world' including the i.e. a copper calorimeter, the surrounding air etc. etc.
Enthalpy H: The heat energy content of a substance. This cannot be determined absolutely but enthalpy changes for a chemical reaction can be measured directly or indirectly from theoretical calculations using known enthalpy values.
Enthalpy change ΔH: The net heat energy transferred to a system from the surroundings or from the surroundings to a system at constant pressure. The Greek letter delta Δ in maths implies a change, in this case a net heat energy change.
A reaction in which heat energy is given out from the system to the surroundings i.e. the enthalpy of the reacting system decreases and the temperature of the system and surroundings rises.
A reaction in which the system takes in or absorbs heat energy from the surroundings i.e. the enthalpy of the system increases and the temperature of the system and surroundings falls OR the system must be heated to initiate the reaction and provide the heat absorbed.
The two diagrams below illustrate how exothermic (left) and endothermic (right) reactions are specified on an enthalpy level diagram.
Standard Enthalpy of Reaction ΔHr/react/reaction is the enthalpy change (heat absorbed/released, endothermic/exothermic) when molar quantities of reactants as stated in an equation react under standard conditions (i.e. 298K/25oC, 1 atm/101kPa)
Standard Enthalpy of Formation ΔHf/form/formation is the enthalpy change when 1 mole of compound is formed from its constituent elements with both the compound and elements in their standard states ('normal stable states) i.e. at 298K/25oC, 1 atm/101kPa
The standard state is the most stable state at the standard temperature and pressure e.g. at 298K/25oC and 1 atm/101kPa
Standard Enthalpy of Combustion ΔHc/comb/combustion is the enthalpy change when 1 mole of a fuel (or any combustible material) is completely burned in oxygen (or air containing oxygen) equated to standard conditions (298K/25oC, 1 atm/101kPa).
Standard enthalpy of neutralisation is the energy released when unit molar quantities of acids and alkalis completely neutralise each other at 298K (pressure effects are insignificant for reactions only involving liquids/solutions/solids)
Bond Enthalpy ('bond energy')
This is the average energy absorbed to break 1 mole of a specified bond when all species involved are in the gaseous state.
e.g. for (i) H2(g) ==> 2H(g) ΔH = +436 kJ mol–1 for the H–H bond
It is always endothermic and the reverse process – bond formation, is always exothermic. In many cases the values are averaged from a variety of 'molecular' situations. More on this in the bond enthalpy section.
Some examples of points made on this page with reference to an enthalpy level change diagram
General points: Arrows pointing downwards represent exothermic changes and arrows pointing upwards represent endothermic changes
1. The energy released when 1 mole of aluminium oxide is formed.
2. The endothermic enthalpy of formation of gold(III) oxide
3. This is a much more complex enthalpy level diagram involving hydrogen, chlorine and hydrogen chloride.
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